w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4

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Decolorizing textile wastewater with Fenton’s reagentelectrogenerated with a solar photovoltaic cell

Sandra Figueroa, Leticia Vazquez, A. Alvarez-Gallegos*

Centro de Investigacion en Ingenierıa y Ciencias Aplicadas, UAEM. Av., Universidad 1001 Col. Chamilpa, Cuernavaca, Morelos,

CP 62209, Mexico

a r t i c l e i n f o

Article history:

Received 13 July 2008

Received in revised form

25 September 2008

Accepted 7 October 2008

Published online 18 October 2008

Keywords:

Fenton’s reagent

Azo-dye degradation

Wastewater treatment

Filter-press reactor

Solar energy

* Corresponding author. Fax: þ52 777 329 708E-mail addresses: sandryjaz@yahoo.co

(A. Alvarez-Gallegos).0043-1354/$ – see front matter ª 2008 Elsevidoi:10.1016/j.watres.2008.10.014

a b s t r a c t

In this work it is demonstrated that Fenton’s reagent can be electroproduced with abun-

dant and cheap feedstock: oxygen saturated wastewater and solar energy. Experiments

were carried out in a divided electrochemical flow cell using two electrodes: a three

dimensional reticulated vitreous carbon cathode and stainless steel anode. Fenton’s

reagent is produced by oxygen reduction on the cathode in the presence of 1 mM Fe2þ. The

influence of electrolyte nature and its concentration and potential difference on the

current efficiency, as well as the rate of Fenton’s reagent electroproduction is discussed

and it is concluded that over this extended range of conditions the current efficiency, for

Fenton’s reagent production, fell within the range 50–70%. It is possible to electroproduce

a stoichiometric amount of Fenton reagent for the oxidation of 0.061 mM Reactive Black 5

(in tap waterþ 0.05 M Na2SO4, zpH 2.8). Similar results were obtained for solutions con-

taining 0.1 mM Acid Green 25. Some practical applications in the field of water treatment

are included. The energy required for drive electrochemical reaction is supplied to the flow

cell by means of a commercial solar panel.

ª 2008 Elsevier Ltd. All rights reserved.

1. Introduction carcinogenic and mutagenic to humans (Chung and Stevens,

The elimination of dyestuff from industrial wastewater is an

important environmental target. It is estimated that more

than 10 000 dyes are consumed in textile processing industries

(Neamtu et al., 2002; Meric et al., 2004) and their concentra-

tions are in the range from 10 to 10 000 mg l (Meric et al., 2004),

depending on the process. Synthetic dyes represent a large

problem because they cannot be destroyed by biological

action. Color is one of the most hated pollutant, because

several reasons: i) it is visible and even small quantities of

dyes (�0.005 mg l) are not allowed (He et al., 2007), ii) color can

interfere with transmission of sunlight into natural streams,

iii) many of the azo dyes and their intermediate products,

such as aromatic amines, are toxic to aquatic life,

4, þ52 777 329 777 984.m.mx (S. Figueroa), l

er Ltd. All rights reserved

1993; Lucas and Peres, 2006). Consequently, dyes have to be

removed from textile wastewater before discharge. In prin-

ciple, the treatment process can be divided in four groups

(Gulyas, 1997): a) separation process, b) degradative process, c)

process that chemically modify wastewater constituents but

do not lead to mineralization and c) preparation process

(addition of certain chemicals for a further treatment) like

breaking of emulsions, flocculation, precipitation, among

others. In recent years, advanced oxidation processes have

been widely investigated in the laboratory. Particularly, the

Fenton’s reagent has been used for treating synthetic indus-

trial effluents and it could be used in real industrial effluents,

but many technical challenges must be overcome before this

approach becomes a realistic wastewater treatment option.

etymagana@prodigy.net.mx (L. Vazquez), aalvarez@uaem.mx

.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4284

Among them it can be mentioned: a) Fenton chemistry, b)

Fenton treatment and c) design, construction and long term

operation of a reliable industrial reactor. These technical

challenges are briefly discussed here and a more realistic

technological option for wastewater treatment is proposed.

1.1. Fenton chemistry

The chemistry behind the Fenton process seems to be very

simple: soluble iron cations (normally FeII/FeIII) interact with

hydrogen peroxide as the primary oxidant. The conversion of

the organic matter is achieved in mild conditions by using the

oxidant that is liberated by a catalytic process, in which

a redox chain is maintained between FeII/FeIII. It has been

recognized (Teel et al., 2001) that one of the attractiveness of

Fenton chemistry approach is its near-stoichiometric gener-

ation of a strong intermediate oxidant, as the responsible for

organic oxidation. In this way, the fraction of the organic

conversion (or the fraction of the hydrogen peroxide

consumed) provides a measure of the efficiency of the

oxidation process. However, it was needed more than one

century to understand that the simple chemical interaction

between H2O2 and FeII/FeIII turns out, surprisingly, in a very

complex chemistry and the identification of the strong inter-

mediate oxidant is still controversial. The first feasible inter-

mediate proposed, as the responsible of the organic oxidation,

was a hydroxyl radical, OH� (Haber and Weiss, 1934). Although

the existence of hydroxyl radical was not settled satisfactorily

it was accepted that the main hydroxyl radicals production

from chemical interaction between H2O2 and Fe2þ/Fe3þ could

be explained according to the following mechanism (Walling,

1975):

Fe2þ þ H2O2 / Fe3þ þ OH� þ HO� (1)

Fe3þ þ H2O2 / Fe2þ þ Hþ þ HO�

2 (2)

The hydroxyl radical hypotheses was reinterpreted by

several research groups (Boye et al., 2003), accepting that OH�

and HO2� were the main oxidizing species in Fenton reaction,

but the interactions of H2O2 with ferrous ions produce soluble

ferric complex:

Fe2þ þ H2O2 / FeðOHÞ2þ þ OH� (3)

FeðOHÞ2þ þ H2O2 4 Fe2þ þ H2O þ HO�

2 (4)

FeðOHÞ2þ þ HO�

2 / Fe2þ þ H2O þ O2 (5)

Fe3þH2O2 4 Fe–OOH2þ þ Hþ (6)

Fe–OOH2þ/ Fe2þ þ HO�

2 (7)

In the presence of organics, Fenton chemistry is even more

complex because hydroxyl radical, both iron cations and the

oxidation products enter into a series of consecutives and

parallels reactions.

Simultaneously, an old proposal (Bray and Gorin, 1932) was

retaken by several research groups (Rahhal and Richter, 1988)

and the high-valent iron-oxo intermediates (i.e. FeIV) were

accepted as the responsible of the substrate oxidation. Under

this approach, both intermediate oxidants (hydroxyl radical

and high-valent iron species) are produced by H2O2 interac-

tions with FeII/FeIII, according to the following consecutive

reactions (Bossmann et al., 1998):

Fe2þ þ H2O2 / FeðH2O2Þ2þ (8)

FeðH2O2Þ2þ/ Fe4þðOH�Þ2 / initiate the oxidation (9)

Fe4þ�OH��

2/Fe3þ

aq ð?Þ þOH� þOH�

/continue the oxidation

(10)

The high-valent iron-oxo intermediates hypotheses were

reinterpreted, based on spectrophotometric measurements of

evolved O2 and the disappearance of Fe2þ, proposing that FeIV

is formed according to (Kremer, 1999):

Fe2þ þH2O2 4k1

k2

�Fe2þH2O2

�/k3

�H2OFeO2þ (11)

At this point, Fenton chemistry turns out very complex. The

FeIV can react either with Fe2þ ions to produce Fe3þ or with

H2O2 to produce O2 and Fe2þ or even with Fe3þ to form

a binuclear species or it can also decompose back into FeIV and

Fe3þ, according to:

FeO2þ þH2O2 /

k4Fe2þ þO2 þH2O

8>>><>>>:

Fe2þ þH2O /k5

2Fe3þ þ 2OH�

Fe3þ 4k6

k8

FeOFe5þ ð12Þ

The mixed valence binuclear species (FeOFe5þ) can react

with H2O2 to form O2 and a mixture of Fe2þ/Fe3þ, according to:

FeOFe5þ þH2O2 /k7

O2 þ Fe2þ þ Fe3þ þH2O (13)

According to experimental results (Kremer, 1999), it is

difficult to support the existence of radical species (OH�,

HO2� ) as the main responsible of Fe2þ and substrates oxida-

tion in the Fenton chemistry. Organic substrate (R–H)

oxidation can be understood by the following reactions. The

active intermediate may oxidize stoichiometrically organic

substrate according to one-equivalent reducing agent:

FeO2þ þ RH / R þ Fe3þ þ OH� (14)

And/or two-equivalent reducing agent:

FeO2þ þ H2R / R þ Fe2þ þ H2O (15)

The controversy is still open, because the free radical (OH�)

existence during Fenton chemistry has been proved by elec-

tron paramagnetic resonance by several research groups

(Rosen et al., 2000). But, it is also proved, as it was documented

above, the existence of high-valent iron-oxo intermediates

(i.e. FeIV). Hence, it is not clear the mechanism that liberates

the enormous oxidant power for organic oxidation in mild

conditions.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 285

1.2. Fenton treatment and reactor configuration

Although the chemistry behind Fenton’s treatment is not well

understood, what is very clear is that a mixture of H2O2/FeII or

H2O2/FeIII produces a strong oxidant capable of oxidize (large

and small) organic pollutants in mild conditions and this

approach can be applied in wastewater treatment. In order to

reflect the oxidizing power of the Fenton chemistry, the

following reduction potentials are given in volts (Bossmann

et al., 1998), with respect to NHE: E0(OH�/H2Oaq)¼ 2.9 V,

E0(Fe3þ/FeO2þ-chelate)¼ 0.9 V and it is accepted that HO2� is

less reactive than OH�. The simplest wastewater treatment by

means of Fenton chemistry, in which diluted H2O2 is added to

a mixture of FeII/substrate solution, is the traditional labora-

tory approach used to assess its performance (Kuo, 1992;

Swaminathan et al., 2003; Meric et al., 2004; Lucas and Peres,

2006). Under this approach, a few large scale treatment

processes have been documented (Oliveros et al., 1997; Bae

et al., 2004). Although in all these cases Fenton chemistry

seems to be a good option to treat synthetic industrial effluent

the main drawback of these approaches is the use of

commercial hydrogen peroxide, including the cost and

hazards associated with the transport and handling of

concentrated hydrogen peroxide. In general, all the reactors

described above are impractical for full-scale (the reactor is

just a flask ranging from 100 ml to 1000 ml) industrial waste-

water treatment.

A better approach is based on an old method for electro-

generation of diluted H2O2: cathodic reduction of dissolved

oxygen (Berl, 1939) in an appropriated material (i.e. carbon),

according the following reaction:

O2 þ 2Hþ þ 2e� / H2O2 (16)

As soon as H2O2 is electroproduced it is activated by soluble

Fe2þ to produce a strong oxidant (see Eq. (11)). In wastewater

treatment, the feasibility of this electrochemical route has

been recognized since almost 25 years (Sudoh et al., 1985;

Oturan et al., 2001; Boye et al., 2006; Guinea et al., 2008).

Although the performance of the Fenton chemistry in waste-

water treatment is promising under this approach, the cost

and durability or structural strength of materials used to

construct the electrodes will be a key factor. Platinum is

usually employed to make anodes while carbon (paper,

graphite bar or carbon felt) is used for cathodes. Pt-anodes are

not suitable for large scale experimentation and, although

carbon is a good cathode material, H2O2 electroproduction is

better carried out on high-surface area electrodes due to the

sluggish kinetics and mass transport restrictions due to the

low solubility of oxygen (Pletcher and Walsh, 1993). Other

important feature is the reactor size (<300 ml) and its config-

uration (undivided glass cell or H-type dual-compartment cell)

because they are impractical for full-scale. A successful scale-

up procedure includes the analysis of a variety of dimen-

sionless groups which describe the geometric, kinematic,

thermal, chemical and electrical characteristics (Frıas-Ferrer

et al., 2008).

Electroproduction of Fenton reagent involves two electrode

processes, resulting in hydrogen peroxide production at the

cathode and evolution of oxygen at anode. An alternative

approach is to change the anodic reaction by a less energy

demanding reaction such as iron oxidation:

Fe–2e� / Fe2þ (17)

The advantage of this approach is that the iron ions are

electrochemically produced in the reactor from a sacrificial

iron anode. This approach overcomes the handle and use of

commercial ferrous salt. Under this approach, the oxidation of

gallic acid (Boye et al., 2006) and herbicides (Boye et al., 2003)

were carried out by Fenton reagent electroproduced in an

undivided cell. Iron ions were electroproduced from an iron

anode while H2O2 was electroproduced from a carbon

cathode. However, the continuous ejection of ferrous ions into

the solution can present some disadvantages: (i) Iron ions

concentrations being high, the coagulation process takes

place. Pollutants are thus precipitated in sludge instead of

oxidation/mineralization. The process is not more ‘‘electro-

Fenton’’, it is called ‘‘peroxi-coagulation’’; (ii) High concen-

tration of Fe(II) ions inhibits the efficiency of Fenton chem-

istry. Using the same electrochemical technique, both

Fenton’s reagent species (Fe2þ/H2O2) were simultaneously

electroproduced at large scale for treating a real wastewater

(Duran Moreno et al., 2004). In spite of the sludge production,

the organic degradation by this method is good. However, the

electrogeneration of Fenton’s reagent under this condition is

facing the challenges related to reactor configuration, design,

construction and long term operation.

In this study it is accepted that chemical interaction

between H2O2 and FeII/FeIII produces a strong intermediate

oxidant (FeO2þ) and it is the responsible for organic oxidation.

Under this approach, it is demonstrated that FeO2þ can be

indirectly electroproduced with abundant and cheap feed-

stock: oxygen saturated industrial effluent and solar energy.

In fact, H2O2 is electroproduced from cathodic O2 reduction in

a divided parallel plate flow-cell reactor. A commercial 304

stainless steel mesh anode is separated from a three dimen-

sional reticulated vitreous carbon cathode by a Nafion 117

membrane. The potential difference between cathode and

anode is applied by means of a photovoltaic panel. Following

the main guidelines in advanced oxidation processes for

wastewater treatment (Gulyas, 1997), this approach can be

used to transform recalcitrant organic wastewater constitu-

ents to biodegradable compounds rather than mineralize

them. The flow-cell is a parallel plate reactor and its configu-

ration is reliable for scaling-up procedures (Walsh, 1993).

2. Theoretical approach

On carbon electrodes oxygen reduction can proceed mainly in

a 2-electron pathway giving hydrogen peroxide in acid media

according to Eq. (16), while on the anode (i.e. platinum) water

oxidation is expected, see Eq. (18):

H2O–2e�/ ð1=2ÞO2 þ 2Hþ (18)

The overall reaction is:

H2O þ ð1=2ÞO2 / H2O2 (19)

To increase the competitiveness of hydrogen peroxide

production, oxygen evolution can be replaced by a more

Fig. 1 – Reactive Black, MW 999.8, l [ 597 nm and Acid

Green 25, MW 622.6, l [ 642 nm.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4286

energetically favorable electrochemical reaction like iron

oxidation, see Eq. (17).

The theoretical energy required for driving electrochemical

reaction (16) is: DG0¼ 237.1 kJ mol�1, while electrochemical

reaction (17) is thermodynamically spontaneous:

DG0¼ 78.9 kJ mol�1. However, under anode potential, Fe2þ can

be oxidized to Fe3þ, according to

Fe2þ–e� / Fe3þ (20)

The theoretical energy for electrochemical reaction (20) is

DG0¼ 74.2 kJ mol�1, that is 162.9 kJ mol�1 less than the elec-

trochemical reaction (18). For a Fe-anode and a carbon

cathode the expected electrochemistry in the flow-cell can be

represented by an anodic reaction (Eq. (20)), a cathodic reac-

tion (Eq. (16)) and the sum of them, the following global

reaction:

2Fe2þ þ O2 þ 2Hþ/ 2Fe3þ þ H2O2 (21)

Electrochemical reaction (19) requires 116.7 kJ mol�1

(ECell0 ¼ 0.624 V), while electrochemical reaction (21) just needs

28 kJ mol�1 (ECell0 ¼ 0.145 V). The electroproduction of iron ions

from iron anode could be taken in pro to form Fenton reagent

(H2O2/Fe2þ) in the catholyte as long as their transfer rate

through the cation permeable membrane is enough. Indeed,

the transfer rate from the anolyte to the catholyte depends on

several factors, among them it can be mentioned cations

concentration in the anolyte and the rate of iron cations

immobilization on the membrane surface. The flow-cell

performance (the optimum voltage, the cell current and

current efficiency for hydrogen peroxide electroproduction,

and hence, FeO2þ production) was experimentally found and

theoretical stoichiometric equations were derived from Eqs.

(14) and (15) for oxidizing known organic substrates. Realistic

theoretical predictions for decolorizing wastewater treatment

are feasible by coupling the flow-cell performance to the

Faraday’s law.

3. Experimental

3.1. Solutions and chemicals

In order to give a more realistic approach, all aqueous solu-

tions were prepared using tap water and its chemical

composition (major cations and anions) is shown in Table 1.

Samples of industrial synthetic dyes such as Reactive Black

5 (RB5) and Acid Green 25 (AG25) were supplied by Ciba

Specialty Chemicals and were used as obtained without

further purification. Fig. 1 shows the structure of these

synthetic dyes. An industrial aqueous sample, coming from

a dying bath effluent, was given by a Mexican textile industry

Table 1 – Chemical composition of the tap water.

Ion Naþ Mg2þ Ca2þ Kþ Cu2þ

ppm 130 5.5 3.2 2.6 <0.02

pH 7.5

for degradation in the flow-cell. The rest of chemicals used in

this work were of reagent grade quality and were used as

obtained from the supplier (Aldrich-Sigma or JB Baker)

without further purification.

3.2. Electrodes and cell

The electrochemical cell is a parallel plate reactor and it is

fully described elsewhere (Alvarez-Gallegos and Pletcher,

1998). In this work a modified version of the flow-cell is

studied and descriptions of main changes are given here. The

cathode was a three dimensional (RVC) electrode, 60 pores per

inch, purchased in bulk form (Electrolytica Inc., NY) and

machined to meet the required size (5 cm� 5 cm� 1 cm). The

specific surface area was 40 cm2 cm�3 of RVC. The superficial

cathode area in the direction of current flow was 25 cm2. The

RVC electrode was glued to stainless steel current collector

surface by means of a silver conductive epoxy (supplied by

Pelco International) and the rest of the surface was insulated

using insulator paint. The anode was a gauze (5 cm� 5 cm) of

commercial 304 stainless steel (304 SS). Catholyte and anolyte

were separated by a cation permeable membrane (Nafion�

117). The electrochemical cell (Fig. 2a) is made of four blocks of

acrylic and sheets of silicon rubber gaskets were collocated

between them to avoid leaks. This cell configuration, allows

two possible anode positions in the anolyte compartment: at

the end-plate (between the blocks of acrylic the anode will be

Fe2þ Cl� SO42� PO4

3� Total alkalinityas CaCO3

0.013 188 2.16 <0.001 65

Fig. 2 – a: Main parts of the flow-cell: four blocks of acrylic, 3D-RVC electrode (5 cm 3 5 cm 3 1 cm), 60 pores per inch, fitted

tightly into the centre of the catholyte channel, cation permeable membrane separator (Nafion� 117), stainless steel gauze

anode (5 cm 3 5 cm) and silicon rubber gaskets. b: Hydraulic circuit, flow-cell and solar panel.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 287

z1 cm far from the cathode, this position is shown in Fig. 2a)

and near the membrane (between two gaskets the anode will

be a few millimeters away far from the cathode).

Electrochemical experiments were controlled by a solar

panel (50 Watts, 17 V, 2.9 A, Model BP 350U) purchased from

Syscom Inc. The applied potential difference (DECell) between

the cathode and anode was selected by using a home built

voltage regulator and the electrical charge was integrated

from tabular data (current vs. time) taken manually from

a voltage regulator display. The hydraulic circuit, the flow-cell

and the solar panel are sketched in Fig. 2b.

All electrolytic experiments were performed in the flow-

cell described above containing 1.5 l of 0.8 M H2SO4 in the

anolyte compartment. The presence and quantity of Hþ in the

anolyte trend to drop the pH in the catholyte due to Hþ

transfer rate through the cation permeable membrane, from

the anolyte to the catholyte. Under this condition it is difficult

to keep an optimum pH of 2.8 in the catholyte. Therefore in

this work it is decided to have a variable volume of catholyte

made of 0.05 M Na2SO4 at pH z 2, adjusted with H2SO4, unless

it is stated otherwise. During electrolysis time oxygen was

bubbled into catholyte. Catholyte and anolyte volumetric

flows rates were both 8 l min. Experiments were carried out at

room temperature.

3.3. Procedures

The quantitative analysis of H2O2 was performed following

the classical permanganate method (Kolthoff and Belcher,

1957). During electrolysis time samples (10 ml) were

0

2

4

6

8

10

12

14

0 1000 2000 3000 4000 5000 6000 7000Charge / C

Hyd

ro

gen

P

ero

xid

e / m

M

Fig. 3 – Hydrogen peroxide production vs. charge passed

during electrolyses of 0.05 M Na2SO4 (pH z 2) continuously

saturated with oxygen. The electrolyses were carried out at

constant potentials of (B) DECell [ 2.0 V, 4 [ 70 ± 5%; (-),

DECell [ 2.5 V, 4 [ 70 ± 5% and ( ) DECell [ 3.0 V,

4 [ 50 ± 5%.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4288

systematically taken and immediately analyzed for hydrogen

peroxide determination. A table was manually formed with

the following values: time (minutes), observed current (A),

volume of 0.1 N KMnO4 used and M H2O2 produced. From this

table, current vs. time were plotted and by numerical inte-

gration the electrical charge was found. The current efficiency

of hydrogen peroxide production was calculated using Far-

aday’s law.

The Chemical Oxygen Demand (COD) was used to estimate

the amount of organic matter in the synthetic and real

wastewater samples and it was analyzed by a HACH� diges-

tion procedure: organic matter is destroyed by a boiling

(150 �C) mixture of chromic and sulfuric acids. A silver

compound is added as a catalyst to promote the oxidation of

organics, and a mercuric compound is included to reduce

interference from the oxidation of chloride ions by dichro-

mate. A sample is refluxed with known amounts of potassium

dichromate and sulfuric acid during 2 h. End products are

carbon dioxide, water and various chromium species. The

excess dichromate is determined by using a HACH� Spectro-

photometer DR/4000 UV–Vis, the limit of detection of this

procedure is 10 ppm COD. Absorbance was measured as

a function of the wave length by means of a HACH� spectro-

photometer DR/4000 UV–Vis. Total iron (Fe2þ and Fe3þ) was

Table 2 – Hydrogen peroxide production in the flow cell as a funin oxygen saturated acid solutions (pH z 2). Applied DECell [ 2

IS (mol l�1) 1.3 l of Catholyte madeof tap water plus sup

de

0.15 0.05 M Na2SO4

0.09 0.03 M Na2SO4

0.03 0.01 M Na2SO4

0.15 0.024 M Na2SO4þ 0.078 NaCl

0.09 0.029 M Na2SO4þ 0.003 NaCl

0.03 0.009 M Na2SO4þ 0.003 NaCl

analyzed by FerroZine method: the FerroZine Iron Reagent

forms a purple-colored complex with iron ions in samples that

are buffered to a pH of 3.5. Ferrous iron was analyzed by 1,10-

phenanthroline method: the 1,10-phenanthroline indicator in

the Ferrous Iron Reagent reacts with ferrous iron in the

sample to form an orange color in proportion to the Fe2þ ion

concentration. Ferric iron does not react. The ferric iron

concentration can be determined by subtracting the ferrous

iron concentration from the results of a total iron test. Both

are spectrophotometric HACH� methods.

4. Results and discussions

4.1. Hydrogen peroxide production

Hydrogen peroxide production was investigated as a function

of three parameters: the applied potential (DECell), supporting

electrolyte and the relative anode position with respect to the

cathode. Experiments were carried out using 1.3 l of catholyte

in the two-electrode configuration electrochemical flow-cell,

with the anode position near the membrane, between two

gaskets, so the anode is placed a few millimeters far away

from the RVC cathode.

When DECell< 1.5 V was applied, little current was

observed but hydrogen peroxide was not detected, probably it

was under the limit of detection of the permanganate method.

When 1.5 V<DECell< 1.7 V, some current (z70 mA) was

observed and a low hydrogen peroxide concentration was

detected. Under these conditions, the total amount of

hydrogen peroxide that was electrogenerated in the catholyte

for about 200 min of electrolysis was 1 mM. By increasing the

applied DECell, hydrogen peroxide starts to accumulate further

in the catholyte. Fig. 3 shows data from three different elec-

trolyses of 200 min each, to reduce O2. It can be seen that

hydrogen peroxide production can accumulate in the cath-

olyte and is a linear function of the electrical charge passed

during the oxygen reduction in 0.05 M Na2SO4 at different

potentials. It can be seen that, in the potential interval of

2.0 V<DECell< 2.5 V, the current efficiency (f) of hydrogen

peroxide production (experimental slope z 70� 5%) is not

significantly affected by the applied potential during the

electrolysis. At DECell¼ 3.0 V, the current efficiency, for the

formation of H2O2, drops to 50� 5%. If the applied potential

ction of ion strength (IS), nature, and concentration of ions.0 V. Electrolysis time: 180 min at room temperature.

Averagederficial currentnsity/(A cm�2)

Averagedcurrent

efficiency

Averagedamount ofH2O2/(mM)

0.0052 70% 6–7

0.0042 50% 3

0.0032 50% 2

0.0048 50% 3

0.0040 50% 2–3

0.0023 50% 1–2

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 289

increases more than DECell¼ 3.0 V, the current efficiency (f)

drops dramatically due to hydrogen evolution reaction.

From Fig. 3 it can be seen that after the passage of 4000 C (at

DECell¼ 2.5 V) 11 mM H2O2 were formed. Similar results were

obtained using a Pt-anode (Alvarez-Gallegos and Pletcher,

1998). Moreover, current efficiencies were similar regardless if

the anode material was platinum or stainless steel. When the

above experiments were repeated, moving the anode position

1 cm away from the cathode (the inter electrode gap changed

from a few millimeters to 1 cm) the main parameters

(observed current, hydrogen peroxide production and current

efficiency) were essentially the same. Similar results were

obtained increasing the anode area (using a stack of a SS

anode meshes). Under these experimental conditions,

hydrogen peroxide production does not depend on both: the

anode position (relative to the cathode) and the number of

anode meshes.

Hydrogen peroxide production was systematically studied

at DECell¼ 2.0 V as a function of ionic strength (IS), the nature,

and concentration of different ions in acidic solutions. Main

results are presented in Table 2.

As it was expected, as the IS decreases, the hydrogen

peroxide production drops markedly, see the trend of the

three first rows of Table 2. However, the presence and

quantity of NaCl trend to drop hydrogen peroxide production

and its current efficiency, see the first and the fourth rows in

Table 2. The effect of NaCl in retarding both H2 evolution and

H2O2 electroproduction was previously documented (Alvarez-

Gallegos and Pletcher, 1998). A possible explanation could

be the diminishing of the cathode electroactive area in the

presence of Cl�. However, further experimental evidences

are needed for making a better conclusion on that point. As

the DECell increases, the current efficiency drops due to

hydrogen evolution and this effect is more marked in

solutions with a lower ionic strength. Tap water is a good

background for hydrogen peroxide electrogeneration as long

as some salt (Na2SO4 rather than NaCl) is added to increase

the ionic strength. Over an extended range of conditions

(i.e. applied potential, anode material, ion strength, inter

electrodes gap, the nature and concentration of major ions)

the current efficiency, for hydrogen peroxide production, fell

within the range 50–70%.

Table 3 – Assessment of iron ions transfer rate through the catithe hydrogen peroxide production in the flow cell. Catholyte wElectrolysis was carried out at DECell [ 2.0 V for 180 min at roo

Electrolysestime/(min)

Catholyte

Fe(Total)ppm

Fe(II)ppm

Fep

0 0 0 0

20 0.7 0.4 0

40 0.53 0.61 0

100 0.54 0.15 0

120 0.75 0.15 0

140 0.69 0.18 0

160 0.82 0.18 0

180 0.81 0.29 0

4.2. Electrogeneration of Fenton reagent(probably FeO2þ)

As it was mentioned above, the electroproduction of iron ions

(Fe2þ/Fe3þ) on the 304 SS anode could be taken in pro to form

Fenton’s reagent (H2O2/Fe2þ) in the catholyte. However, under

the experimental conditions studied here it was not possible

because the transfer rate of iron ions through the cation

permeable membrane was not enough to form it. The iron

ions (Fe2þ/Fe3þ) transfer rate was assessed by determining

total iron, ferric iron ion and ferrous ion in both solutions:

anolyte and catholyte. Table 3 shows data from electrolysis to

reduce O2 on a 60 ppi RVC cathode in the flow-cell reactor.

Catholyte was 1.3 l of 0.078 M NaClþ 0.024 M NaSO4. Electrol-

ysis was carried out at DECell¼ 2.0 V for 180 min at room

temperature. During electrolysis time some samples were

systematically taken from both, catholyte and anolyte, and

immediately analyzed for total iron, ferrous ion and hydrogen

peroxide. At the end of the electrolysis z3 mM H2O2 was

accumulated. At the end of the experiment, the accumulation

of Fe3þ in the anolyte was z1.7 mM and Fe2þwas not detected.

Due to the low transfer rate of Fe3þ through the cation

permeable membrane, iron ions concentrations were

0.005 mM Fe2þ and 0.009 mM Fe3þ in the catholyte.

The presence of Fe2þ in the catholyte can be explained by

the Fe3þ reduction on the cathode surface. Under these

experimental conditions, Fenton’s reagent could not be

formed, because in order to form it z1 mM Fe2þ is required.

However, the amount of hydrogen peroxide produced in the

catholyte, under the experimental conditions studied above,

can be used in wastewater treatment if it is activated with

z1 mM Fe2þ. But the main drawback of this approach is the

critical catalytic activation time of H2O2 by iron ions. Indeed,

the couple Fe2þ/Fe3þ regeneration, during the Fenton chem-

istry is not efficient due to chemical speciation partially

described by Eqs. (3)–(13).

Fig. 4 shows the data from two electrolyses carried out in

the flow-cell, at DECell¼ 2.0 V in a catholyte consisting of

0.05 M Na2SO4. In absence of Fe2þ, see curve (a); hydrogen

peroxide production can accumulate in the catholyte and is

a linear function of the electrical charge passed during the

oxygen reduction. In contrast, when 1 mM Fe2þ is added to the

catholyte, the current efficiency of hydrogen peroxide

on permeable membrane, from anolyte to catholyte duringas 1.3 l of 0.078 M NaCl D 0.024 M Na2SO4 at pH z 2.

m temperature.

Anolyte

(III)pm

Fe(Total)ppm

Fe(II)ppm

Fe(III)ppm

0 0 0

.3 0.65 0 0.65

1.9 0 1.9

.39 5.84 0 5.84

.6 60 0 60

.51 69.1 0 69.1

.64 88.3 0 88.3

.52 94 0 94

0

1

2

3

4

5

6

7

0 1000 2000 3000 4000 5000 6000 7000Charge / C

Hyd

ro

gen

P

ero

xid

e / m

M (a)

(b)

Fig. 4 – Electrolyses of 0.05 M Na2SO4 adjusted to pH 2

continuously saturated with O2 at a reticulated vitreous

carbon cathode in a flow-cell. DECell [ 2.0 V. Plots of H2O2

formed vs. charge passed for solution with a) 0 mM added

Fe(II) and b)1 mM added Fe(II).

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4290

production is close to 0%. This suggests that a strong oxidant

or a mixture of them are formed (FeIV, OH� or HO2�) when the

hydrogen peroxide produced reacts with iron in the bulk

solution during the passage of the first 2500–3500 C of charge

(z160–200 min). Thereafter, when the catalytic activity is lost,

the hydrogen peroxide concentration in the catholyte began to

rise again and Fenton chemistry stops, see curve (b). A similar

situation was documented before (Alvarez-Gallegos and

Pletcher, 1998).

The following section describes an application of hydrogen

peroxide electrogenerated with a solar photovoltaic cell for

decolorizing synthetic wastewater containing Reactive Black 5

(RB5) and Acid Green 25 (AG25). A sample of a typical Mexican

textile industrial effluent, containing RB 5 is included as well.

4.3. Reactive Black 5

The complete oxidation of RB5 (C26H21O19S6N5Na4) is a 152

electrons oxidation reaction; if sulfur and nitrogen are trans-

formed in sulfuric and nitric acids see the following equation:

C26H21O19S6N5Na4 þ 72H2O–152e� / 26CO2 þ 152Hþ

þ 4H2SO4 þ 5HNO3 þ 2Na2SO4 (22)

If it is accepted the stoichiometric generation of a strong

intermediate oxidant described by Eq. (11), hence, the

conversion of 1 mol of RB5 to CO2 requires up to 76 mol of

FeO2þ, if Eqs. (11) and (22) are 100% efficient. The oxidation of

RB5 in the flow-cell by means of electrogenerated FeO2þ was

investigated. The theoretical charge and amount of FeO2þ (or

H2O2) for complete oxidation of the RB5 can be evaluated by

means of Faraday’s law and the flow-cell performance (2.0 V,

0.130 A and f¼ 70%):

q ¼ mnFf

(23)

where q is the theoretical charge (in C) passed through the cell,

m is the amount (in mol) of organic matter (or hydrogen

peroxide electrogenerated), n is the number of electrons

involved in the oxidation reaction, F is de Faraday constant

(96 485 C mol�1) and f is the current efficiency for hydrogen

peroxide production (z70%). Substituting the appropriate

values in Eq. (23) we obtain the theoretical charge for complete

oxidation of 1.3 l 0.061 mM RB5 (0.079 mmol RB5):

q ¼ð0:0000793 molÞð152e�Þ

�96485 C mol�1

�

0:70¼ 1661 C (24)

Taking into account the theoretical charge and Faraday’s

law, the theoretical amount of hydrogen peroxide (or FeO2þ)

needed is: 6 mmol if Eqs. (11), (15) and (22) are 100% efficient.

From the Faraday’s law and flow-cell performance, the theo-

retical time (t) necessary for electrogenerating 6 mmol of H2O2

can be estimated: t¼ 1661 C/0.130 A¼ 12 777 s (213 min).

Although the complete oxidation of RB5 is theoretically

feasible, the required time for FeO2þ electrogeneration is at

the limit of the critical time (see Fig. 4) at which Fenton

chemistry stops. However, decolorizing dye wastewater

requires only a small charge, because the color of the dyestuff

is a function of the conjugated double bonds length in the

aromatic molecule. The ring opening process should lead to

discoloration early in the oxidation process.

Fig. 5a reports spectra recorded at various stages in the

electrolysis of an oxygen saturated catholyteþ 0.061 mM

RB5þ 1 mM FeSO4 and pH 2 at DECell¼ 2.0 V. The initial solu-

tion is deep blue (52 ppm of experimental COD) and has to be

diluted by a factor of four before a spectrum, curve (a), can be

recorded; it shows a peak (absorbance¼ 1.167) in the spectrum

at lmax¼ 598 nm. After the passage of only 362 C (45 min), the

color has changed to a yellowish brown and the COD drops to

the detection limit (10 ppm). On continuing the electrolysis, by

the passage of 620 C (75 min), the solution has become pale

yellow (COD< 10 ppm) and the spectrum, curve (c) shows only

an absorption tail into the UV. By 1256 C (140 min), curve (d),

the solution is effectively colorless and the COD was abated

90% from its original value (COD< 10 ppm).

4.4. Acid Green 25

Following the same assumptions as before, the complete

oxidation of AG25 (C28H20O8N2Na2S2) is a 124 electrons

oxidation and the stoichiometric conversion of 1 mol of AG25

to CO2 requires up to 62 mol of FeO2þ, if Eqs. (11) and (22) are

100% efficient. The oxidation of 0.1 mM AG25 in the flow-cell

by means of electrogenerated Fenton’s reagent was investi-

gated. The theoretical charge and amount of FeO2þ for

complete oxidation of 1.5 l 0.1 mM AG25 (0.15 mmol AG25) in

solution are 2894 C and 10.5 mM FeO2þ, respectively. The

theoretical time needed for electrogenerating 10.5 mM of H2O2

are 22 262 s (371 min). Theoretical time is beyond the critical

time at which Fenton chemistry stops. However, color abate-

ment is achieved in a shorter time.

Fig. 5b reports spectra recorded at various stages in an

electrolysis of an oxygen saturated catholyteþ 0.1 mM

AG25þ 1 mM FeSO4 and pH 2 at DECell¼ 2.0 V. The initial

solution is clear green (62 ppm of experimental COD) and the

spectrum is shown in curve (a). The spectra shows two peaks

at l1¼ 643 nm (absorbance¼ 0.916) and l2¼ 610 nm

(absorbance¼ 0.925). After the passage of only 55 C (40 min

and 46 ppm COD), the color has change to a pale green, see

curve (b). On continuing the electrolysis, by the passage of

0.3

0.6

0.9

1.2

1.5

Ab

so

rb

an

ce

diluted x 3

diluted x 2

(a)(b)

(c)

(d)

(e)

0

0.5

1

1.5

2

2.5

350 450 550 650 750 850 950 1050Wavelength / nm

Ab

so

rb

an

ce

(a) diluted x 4

(b)

a

b

(c)

(d)

0

0.5

1

1.5

2

2.5

350 450 550 650 750 850Wavelength / nm

Ab

so

rb

an

ce

(a)

(b)

(c)

(d)(e)

Fig. 5 – a. Spectra as a function of charge passed for

a solution (1.3 l) containing 0.061 mM reactive black 5 in

0.05 M Na2SO4 D 1 mM Fe2D; (a) deep blue, 0 C, 52 ppm

COD; (b) yellowish brown, 362 C, <10 ppm COD; (c) pale

yellow, 620 C, <10 ppm COD; and (d) colorless, 1256 C,

<10 ppm COD. Theoretical charge for complete oxidation

1661 C. Applied DECell [ 2.0 V and ICell [ 0.130 A. b. Spectra

as a function of charge passed for a solution (1.5 l)

containing 0.1 mM acid green 25 in 0.05 M Na2SO4 D 1 mM

Fe2D. (a) Clear green, OC, 62 ppm COD; (b) pale green, 558 C,

46 ppm COD; (c) very pale green, 1176 C, 29 ppm COD; and

(d) colorless, 1836 C, 13 ppm COD. Theoretical charge for

complete oxidation 2894 C. Applied DECell [ 2.0 V and

ICell [ 0.130 A.

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 291

1176 C (80 min, 29 ppm COD), the solution has become very

pale green, see curve (c). By 1836 C (120 min, 13 ppm COD) the

spectrum, curve (d), shows only an absorption tail into the UV

and the solution is effectively colorless. On continuing the

electrolysis the COD drops to the detection limit (10 ppm) by

2160 C (140 min).

0300 400 500 600 700 800

Wavelength / nm

Fig. 6 – Spectra as a function of charge passed for a solution

(1.5 l) containing an industrial effluent sample with an

unknown amount of reactive black 5 in 0.05 M

Na2SO4 D 1 mM Fe2D. (a) Deep dark, 0 C, 230 ppm COD; (b)

deep dark, 202 C, 186 ppm COD; (c) deep dark, 639 C,

179 ppm COD; (d) yellowish brown, 2823 C,

COD [ 143 ppm; and (e) pale brown, 5782 C, 121 ppm COD.

Applied DECell [ 2.5 V and ICell [ 0.200 A.

4.5. Industrial effluent

The oxidation of an industrial aqueous sample, coming from

a dying bath effluent, was investigated following the same

procedure discussed before. The aqueous sample was given by

a Mexican textile industry without details about the process of

dying (i.e. mordant solutions, fixing mordents, rinsing solu-

tions, etc.). The only information given by the textile industry

is that tap water was used to prepare solutions and industrial

RB5 (supplied by Ciba Specialty Chemicals) is used to dye

cottons. This industrial aqueous sample is considered as the

representative textile effluent which is discharged into the

environmental without further treatment.

The industrial aqueous sample was just filtered when

received, it was very dark and, apparently, without oil and

grease. In order to assess its pollution magnitude, several COD

analyses were performed to the sample. The averaged COD

was 2720 ppm, indicating a heavy polluted textile industrial

effluent. A ten-fold dilution was made with tap water and it

was conditioned to have the following sample: 1.5 l of 0.05 M

Na2SO4þ diluted industrial sampleþ 1 mM FeSO4. The oxida-

tion of such industrial sample (230 ppm of COD) in the flow-

cell by means of electrogenerated Fenton’s reagent was

investigated at DECell¼ 2.5 V at room temperature. Theoretical

calculations are more difficult than before because the

chemical details of the effluent were not given. However,

based in the sample COD a rough estimation of the RB5

concentration was made, supposing that only the RB5 is

contributing to the experimental COD. Therefore, the

concentration of RB5 in the diluted industrial sample was

0.27 mM. This concentration is more than 4 times bigger than

the previous one.

The theoretical charge and amount of FeO2þ for complete

oxidation of the diluted industrial sample (taking into account

the new experimental conditions: 2.5 V, 0.200 A and 4¼ 65%)

are 9138 C and 30.8 mM FeO2þ respectively. The theoretical

time needed for electrogenerating 30.8 mM of H2O2 is

t¼ 45 690 s (12.7 h). Theoretical time is, of course, beyond the

critical time at which Fenton chemistry stops. However, color

abatement is achieved in a shorter time. Fig. 6 reports spectra

recorded at various stages in electrolysis of the previous

solution. Although the original solution was diluted, the initial

solution was deep dark and has to be diluted by a factor of

three before a spectrum, curve (a), can be recorded; it shows

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4292

a peak (absorbance¼ 0.97) in the spectrum at lmax¼ 567 nm.

After the passage of 202 C (20 min), solution was still deep

dark (COD¼ 186 ppm) and has to be diluted by a factor of two

before a spectrum, curve (b), can be recorded; the peak in the

spectra has changed from lmax¼ 567 nm to lmax¼ 479 nm; by

the passage of 639 C (60 min) solution was still dark

(COD¼ 179 ppm) and has to be diluted by a factor of two

before a spectrum, curve (c), can be recorded. On continuing

the electrolysis, by the passage of 2823 C (260 min), color has

change to a yellowish brown (COD¼ 143 ppm), see curve (d);

by the passage of 5782 C (450 min), the solution has become

pale brown (COD¼ 121 ppm) and the spectrum, curve (e)

shows only an absorption tail into the UV. Two hours more of

electrolysis was carried out but both parameters color spectra

and COD were essentially the same and electrolysis has to be

stopped. At this point, it is clear that RB5 molecule is broken

and smaller and more oxidized organic molecules are being

formed in solution. Beyond this time the decay in COD do not

changes considerably and it was not clear whether this

decrease in rate results from the more difficulty in oxidizing

the smaller molecules (the presence of very stable by-prod-

ucts) or it occurs because of a change in speciation of iron ions

(Fe2þ/Fe3þ).

Although the oxidation path during the industrial sample

electrolyses is unknown it can be taken, as an approximation,

some paths proposed during biologic degradation of RB5. The

first intermediates due to biological degradation for BR5 are

naphthalenic and benzenic ring amines (Storm, 2002). It is

logic to expect them as the first intermediate species during

the attack of FeO2þ on RB5 because similar organic structures

have been found during sonolysis and ozonation of azo dyes

(He et al., 2007). Most of the naphthalenic ring amines

produced during dye degradation are not commercially

available and they could not be tested but, they should be

oxidized by Fenton chemistry with the same degree of diffi-

culty than the previous dyestuff studied. As an approxima-

tion, aniline can be taken as an organic model representing

benzenic ring amines, cresol and phenol could represent

benzenic ring structures, and oxalic acid could represent

aliphatic acids, during RB5 degradation. The feasibility of

Table 4 – Electrolysis performed in a divided flow-cell (60 ppi 3working at 70% current efficiency for hydrogen peroxide produmolecule D 1 mM Fe(II) at pH z 2. Ph [ phenol; Cr [ cresol; AnpBQ [ parabenzoquinone; OxAc [ oxalic acid; eL [ electron toconcentration in mM; ICell [ averaged cell current in amperes; qcomplete organic molecule oxidation in C; TET [ theoretical eminutes.

mM e� qT/(C) TET(min)

0.37 Ph 28 2399 105

0.33 HQ 26 2129 100

0.33 Cat 26 2129 100

0.33 An 36 2947 123

0.33 pBQ 24 1528 73

0.33 Cr 34 2165 108

10 mM OxAc 2 2757 92

these organics oxidation by means of Fenton chemistry were

tested before in synthetic, acidic wastewaters by Fenton’s

reagent electrogenerated at a reticulated vitreous carbon

cathode using the flow-cell described above (Alvarez-Gallegos

and Pletcher, 1999). The organic molecules considered were:

phenol (Ph), cresol (Cr), aniline (An); hydroquinone (HQ),

catechol (Cat), parabenzoquinone (pBQ) and oxalic acid

(OxAc). Their initial concentrations, main experimental

conditions, theoretical and experimental charge and time for

complete oxidation are shown in Table 4.

According to the theoretical calculations shown in Table 4,

the COD will drop rapidly to below 10 ppm during the elec-

trolysis time for the aniline, oxalic acid and all benzenic ring

structures. Indeed, for these aqueous solutions, the flow-cell

is able to produce the required amount of a strong oxidant

(FeO2þ) before the critical time is reached (160–200 min).

Experimental evidences shown that all benzenic ring struc-

tures (including aniline) were readily oxidized by hydrogen

peroxide in the presence of Fe(II) at room temperature. The

hardest electrolysis was the biodegradable 10 mM OxAc

aqueous solution. In this case, experimental electrolysis was

stop by 300 min (z8600 C) and the COD dropped to 12 ppm.

Hence, the formation of very stable by-products during RB5

oxidation, as it was documented elsewhere (Lucas and Peres,

2006; Swaminathan et al., 2003), cannot be correct and

experimental evidences points towards another feasible

interpretation: the couple Fe2þ/Fe3þ regeneration, during the

Fenton chemistry is not efficient due to chemical speciation.

The energy consumption (EW, in kW h m�3) for wastewater

treatment can be calculated from Eq. (25), taking into account

the electrolysis time (telectrolysis, in hours), assuming an

average cell current (ICell, in Amperes), cell voltage (ECell, in

Volts) and the volume of catholyte (VCatholyte, in m3), main

results are summarized in Table 5.

EW ¼ðECellÞðICellÞ

�telectrolysis

�VCatholyte

(25)

This approach could be attractive for effluent treatment

because the energy consumption for the removal of color in

industrial effluents is not high. Moreover, in future works, the

D-RVC cathode, 5 cm 3 5 cm 3 1 cm, and Pt mesh anode)ction. 1.8 l of catholyte: 0.05 M Na2SO4 D organic[ aniline; HQ [ hydroquinone; Cat [ catechol;be lost by mol of substrate; mM [ Initial organic molecule

T and qE [ theoretical and experimental charge needed forlectrolysis time required for electrogenerate ferryl ion, in

Averagedsuperficial

currentdensity/(A cm�2)

qE/C COD0

(ppm)

0.0152 3100 82

0.0142 3700 74

0.0160 2750 67

0.0192 2100 70

0.0180 2700 55

0.0172 4100 85

0.0160 8600 160

Table 5 – Electrolysis energy consumption for 1 m3 of 0.05 M Na2SO4 D organic aqueous solutions, pH 2. Divided flow-cellwith reticulated vitreous carbon cathode (5 cm 3 5 cm 3 1 cm). Catholyte flow rate 8 l minL1 and room temperature.

1 m3 ofWastewater

DECell (V) ICell/A InitialCOD

(ppm)

Electrolysistime

% ofCOD

depletion

Energy(EW)

(kW h m�3)

0.061 mM

RB 5

2.0 0.200 115 2.33 h

Colorless

solution

>90 0.717

0.1 mM

AG 25

2.0 0.248 62 2.33 h

Colorless

solution

>80 0.770

Industrial

sample

(RB 5)

2.5 0.212 230 7.5 h

Pale brown

z50 2.65

w a t e r r e s e a r c h 4 3 ( 2 0 0 9 ) 2 8 3 – 2 9 4 293

hydrogen peroxide production could be boosted by the use of

a cationic surfactant in the catholyte (Gyenge and Oloman,

2005), improving its current efficiency. If the anolyte is

a closed hydraulic loop, the chromium content (coming from

the stainless steel) should not be an environmental problem. If

anode material is changed (iron anode instead of stainless

steel anode), the accumulation of iron ions (Fe2þ/Fe3þ) in the

anolyte is a problem for a long term operation of a reliable

industrial reactor. Indeed, SEM images showed that, the

Nafion membrane became clog with both Fe(II/III) after 15 h of

reactor operation using an iron anode. However, using

a stainless steel anode the problem is minimized and the

Nafion membrane can last for more than 200 h of reactor

operation. Under these experimental conditions a good option

is to keep the membrane, the stainless steel anode and

withdraw continuously a volume from the anolyte and

replace it with fresh water. This operation will maintain a low

iron ions concentration in the anolyte. Additionally, the iron

ions can be recovered from the withdrawn volume before the

final discharge.

5. Conclusions

It has been demonstrated that Fenton’s reagent (probably

FeO2þ) can be indirectly electroproduced in a flow-cell by

a cathodic reduction of dissolved oxygen on a RVC surface in

the presence of 1 mM Fe2þ, using abundant and cheap feed-

stock: oxygen saturated industrial effluent and solar energy.

The near-stoichiometric generation of a strong oxidant

(probably FeO2þ) can be applied for treating a real textile

effluent as long as the amount of organic matter could be

destroyed before the iron catalytic activity is lost (i.e. 160–

200 min of electrolysis). Experimental conditions were close to

that found in wastewater treatment: tap water (major cations

present, i.e. Naþ, Cl�, SO42�, Fe2þ), pH 2 (to keep iron in solu-

tion), typical azo-dyes concentrations (in range from 10 to

10 000 ppm) and ambient temperature. This approach can be

improved by changing the anode material: stainless steel for

iron. However, for long term operation, the rate of iron cations

immobilization on the membrane surface must be evaluated.

Additionally, organic oxidation may be boosted by replacing

the homogeneous hydrogen peroxide activation for hetero-

geneous catalysis activation.

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